Chemistry 102 Summary July 28th
REDOX (Oxidation-Reduction Reactions)
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Reactions involving the transfer of electrons from one species to
another.
Oxidation: Atom loses electrons. Oxidation number increases.
Reduction: Atom gains electrons. Oxidation number decreases.
Oxidizing Agent: Species that causes oxidation to occur; always the
compound/ion that contains the atom that is reduced.
Reducing Agent: Species that causes reduction to occur; always the
compound/ion that contains the atom that is oxidized.
O(xidation)I(nvolves)L(oss of electrons) R(eduction)I(nvolves)G(ain of
electrons).
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Redox reactions play an important role in everyday life: combustion
of fuels, batteries, corrosion, metal plating of jewelry, industrial
production of chemicals.
Assigning Oxidation Numbers
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Provides a way to keep track of electrons in redox reactions.
For covalent compounds, the oxidation numbers do not represent
the actual charges for the elements in compounds.
Oxidation numbers are imaginary numbers, and are used to
recognize redox reactions.
General Rules
1. The sum of the oxidation numbers for all atoms in a compound/ion
must equal the charge on the compound/ion.
2. Any element by itself has an oxidation number = 0.
Ionic Compounds
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Oxidation number for monoatomic ions equals the charge on the
ion.
Covalent Compounds
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Assign Hydrogen a +1 value.
Assign Oxygen a -2 value.
Assign Fluorine a -1 value.
Use these rules and rule 1 to assign others.
Chapter 4 #67
(a) KMnO4
(b) NiO2
(c) Na4Fe(OH)6
(d) (NH4)2HPO4
(e) Fe3O4
Redox Reactions
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If a reaction is a redox reaction then the oxidation numbers will
change for some of the elements in the reaction.
If there is no change in oxidation numbers it is not a redox reaction.
Chapter 4 #71
(a)
Cu (s) + 2 Ag+ (aq)
(b)
HCl (g) + NH3 (g)
(c)
SiCl4 (l) + 2 H2O (l)
4 HCl (aq) + SiO2 (s)
(d)
SiCl4 (l) + 2 Mg (s)
2 MgCl2 (s) + Si (s)
2 Ag (s) + Cu2+ (aq)
NH4Cl (s)
Balancing Redox Reactions
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Some redox reactions can be balanced by inspection: combustion
and single displacement reactions.
Example #94 Chapter 4
Balance the following reactions by inspection:
(a)
Al (s) + HCl (aq)
(b)
CH4 (g) + S (s)
CS2 (l) + H2S (g)
(c)
C3H8 (g) + O2 (g)
CO2 (g) + H2O (l)
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AlCl3 (aq) + H2 (g)
Most redox reactions are very complicated and are too difficult to
balance by inspection.
Balancing Redox Reactions – The Half-Reaction Method
1. Write the equations for the oxidation and reduction half-reactions.
Sometimes the half-reactions are obvious. However, when in doubt,
assign oxidation numbers to deduce what is oxidized and what is
reduced. Only include the compound/ions which contain the
species oxidized or reduced in the initial half-reactions.
2. For each half reaction:
(a) Balance all elements except hydrogen and oxygen.
(b) Balance oxygen using H2O.
(c) Balance hydrogen using H+.
(d) Balance charge using electrons.
Each half reaction should now be mass balanced and charge
balanced.
3. If necessary, multiply one or both balanced half-reactions by an
integer to equalize the number of electrons transferred in the two
half-reactions (electrons must cancel).
4. Add the half-reactions together, and cancel identical species.
5. Check that the overall reaction is mass and charge balanced.
Balancing Redox Reactions in Basic Solution
After step 4 above (mass and charge balanced equation)
6. To both sides of the equation, add a number of OH- ions equal to
the number of H+ ions. H+ will react with OH- to form H2O.
7. Form H2O on the side containing both H+ and OH- ions, then cancel
H2O molecules appearing on both sides of equation.
8. Check mass and charge balance.
Examples #74 Chapter 4
(a)
(b)
(c)
Cu (s) + NO3- (aq)
Cr2O72- (aq) + Cl- (aq)
Pb (s) + PbO2 (s) + H2SO4 (aq)
Cu2+ (aq) + NO (g)
Cr3+ (aq) + Cl2 (g)
PbSO4 (s)
#75 Chapter 4 Balance the following redox reactions in basic solution
(a)
Al (s) + MnO4- (aq)
(b)
Cl2 (g)
(c)
NO2- (aq) + Al (s)
MnO2 (s) + Al(OH)4- (aq)
Cl- (aq) + OCl- (aq)
NH3 (g) + AlO2- (aq)
THE END
Chemistry 102 Course Summary
Hour Exam 1:
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The Scientific Method
Definition of a Theory and a Law
Accuracy vs. Precision
Random vs. Systematic Error
Significant Figures
Dimensional Analysis
Density
Classification of Matter
Dalton’s Atomic Theory
Nuclear Atom
Atomic Number, Mass Number, Isotope
Metals vs. Nonmetals, Periods, Groups, Ionic vs. Covalent
Nomenclature
Atomic Mass, Avogadro’s number and the mole
Molar Mass, mass %
Empirical vs. Molecular Formula
Chemical Equations, Balancing chemical equations
Stoichiometry and limiting reagent problems
Solvent and solute definitions
Solubility rules
Electrolytes (strong vs. weak vs. non electrolytes)
Concentration and dilution
Precipitation reactions and solution stoichiometry
Definition of acids and bases
Complete vs. Net Ionic Equation
Acid base titrations, equivalence point
Titration type problems
Enthalpy (Endothermic vs. Exothermic process)
Properties of Gases, definition of pressure
Ideal Gas Law
Boyle’s Law, Charles’ Law, Avogadro’s Law
Dalton’s Law of Partial Pressures
Kinetic Molecular Theory
Graham’s Law of Effusion
Real Gas Behavior
Hour Exam II
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Properties of Light (wavelength, frequency)
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Energy associated with electromagnetic radiation
Emission spectra
Bohr model
Ionization Energy
Wave or Quantum Mechanical Model (contributions by Bohr,
deBroglie, Heisenberg, Schrodinger)
Atomic Orbitals (s, p, d and f)
Electron Configurations (Hund’s rule, Aufbau principle) and
exceptions
Valence electrons
Periodic properties and trends (radii trend, ionization energy trend,
electron affinity) and isoelectronic series
Chemical bonding (ionic vs. covalent)
Electronegativity (polar and nonpolar covalent bonds)
Lewis Structures (covalency rules, resonance)
VSEPR (geometry and shape) and predicting molecular polarity
Hybridization (sp, sp2, sp3 and sigma vs. pi bonds)
Intramolecular vs. Intermolecular Forces
Molar heat of vaporization, fusion and sublimation and vapor
pressure
Types of intermolecular forces (electrostatic, dipole-dipole,
hydrogen bonding, London dispersion)
Equilibrium, equilibrium position, K (significance of the value of K)
KP and how KP relates to K
Manipulating K and heterogenous equilibrium
Hour Exam III
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ICE table
Reaction Quotient (Q)
Le Chatelier’s principle (Effect of temperature on value of K)
Acids and Bases (strong vs. weak), Ka, Kb, conjugate acid base
pairs, KW, pH and pOH
pH calculations of strong acids and bases in water
pH calculations of weak acids and weak bases in water
Step-wise Ka expressions for polyprotic acids
Acid base properties of metal and nonmetal oxides in solution
Buffers (Henderson-Hasselbach equation)
Titrations (titrant, equivalence point, neutralization reaction, titration
curve)
Strong acid by strong base and strong base by strong acid titrations
Weak acid by strong base and weak base by strong acid titrations
Polyprotic acid titrations
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Redox reactions (assigning oxidation numbers, balancing redox
reactions)